1. REDOX REACTIONS

2. REDUCTION Previously: What happened to oxygen when it reacted – During reactions oxygen would take on electrons Now: When any element gains electrons

3. OXIDATION Previous: What happened to an element when it reacted with oxygen – During reaction, oxygen would take the electrons from elements Now: When any element loses electrons

4. REDOX Reduction and oxidation NEVER happen by themselves – When an element is oxidized, there is another element that is reduced – When an element is reduced, there is an another element that is oxidized LEO the Lion says GER – Lose an electron  oxidized – Gain an electrons  reduced

5. OXIDATION NUMBERS Oxidation numbers can be assigned using the periodic table if the compound is an ionic compound. Covalent compounds are made of two nonmetals, which from the periodic table are always expected to be negative. – Since covalent compounds are neutral, it is not possible for every element to retain its negative oxidation number. – Only the more electronegative element stays negative; least electronegative element changes to positive Oxidation number is different from formal charge. Using oxidation number gives us another way to account for electrons in chemical changes.

6. RULES FOR DETERMINING OXIDATION NUMBERS 1.The oxidation number of any uncombined element is zero 2.The oxidation number of a monatomic ion equals its charge 3.Oxygen’s oxidation number is -2, except in peroxides (H 2 O 2 ) where it is -1 or when it bonds with fluorine where it will be +2 4.The oxidation number of hydrogen is +1 except when it bonds with metals to from metal hydrides, where it is -1 or in the polyatomic ion NH 4 where it is also -1 5.The sum of the oxidation numbers for a compound must equal zero 6.The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its charge.

7. RULES FOR DETERMINING OXIDATION NUMBERS 1-2: 2Na + Cl 2  2NaCl 3-4: H 2 O 5: CaCl 2 Ca(OH) 2 6: NO 3 - SO 4 -2

8. REDOX REACTIONS Any reaction where the oxidation number of an element is different on the two sides of the chemical equation is a redox equation Mg + 2HCl  MgCl 2 + H 2 Oxidation = loss of electrons Reduction = gaining electrons You will represent redox reactions using half reactions

9. REDOX REACTIONS Half Reactions – Mg + 2HCl  MgCl 2 + H 2 The total exchange of all electrons is accounted for in the combined half reactions

10. AGENTS Reducing Agent: An element that serves as the source of electrons is known as the reducing agent – The reducing agent is the one that is oxidized Oxidizing Agent: An element that receives the electrons is known as the oxidizing agent – The oxidizing agent is the one that is reduced

11. EXAMPLES 4Fe + 3O 2  2Fe 2 O 3 2Fe 2 O 3  4Fe + 3O 2 AgNO 3 + Mg  Mg(NO 3 ) 2 + Ag

12. REDOX REAL-LIFE EXAMPLES Galvanic Wet Cell

13. REDOX REAL-LIFE EXAMPLES How does this work – If you bathe two different strips in a conductive solution, and connect them externally with a wire – Reactions between the electrodes and the solution furnish the circuit with charges continually. – The process that produces the electrical energy continues and becomes useful. – Spontaneously conversion of chemical energy to electrical energy. – The Copper (Cu) atoms attract electrons more than do the Zinc (Zn) atoms. – Zinc is more active and yields its electrons more easily.

14. REDOX REAL-LIFE EXAMPLES

15. AN ALKALINE BATTERY – Anode: Zn cap: Zn(s) → Zn2+(aq) + 2e- – Cathode: MnO2, NH4Cl and carbon paste: 2 NH4 +(aq) + 2 MnO2(s) + 2e- → Mn2O3(s) + 2NH3(aq) + 2H2O(l) – Graphite rod in the center - inert cathode. – Alkaline battery, NH4Cl is replaced with KOH. – Anode: Zn powder mixed in a gel:

16. REDOX REAL-LIFE EXAMPLES

17. CORROSION OF IRON – Since E°(Fe2+/Fe) < E°(O2/H2O) iron can be oxidized by oxygen. – Cathode O 2 (g) + 4H+(aq) + 4e- → 2H 2 O(l). – Anode Fe(s) → Fe 2+ (aq) + 2e-. – Fe 2+ initially formed – further oxidized to Fe 3+ which forms rust, Fe 2 O 3 xH 2 O(s).